The Haber Process for the manufacture of ammonia THE HABER PROCESS This page describes the Haber Process for the manufacture of ammonia from nitrogen and hydrogen, and then goes on to explain the reasons for the conditions used in the process. It looks at the effect of temperature, pressure and catalyst on the composition of the equilibrium mixture, the rate of the reaction and the economics of the process. Important: If you aren't sure about using Le Chatelier's Principle or about the effect of changing conditions on rates of reaction you should explore these links before you go on. When you are reading this page, if you find that you aren't understanding the effect of changing one of the conditions on the position of equilibrium or on the rate of the reaction, come back and follow up these links. A brief summary of the Haber Process The Haber Process combines nitrogen from the air with hydrogen derived mainly from natural gas (methane) into ammonia. The reaction is reversible and the production of ammonia is exothermic. A flow scheme for the Haber Process looks like this: Some notes on the conditions The catalyst The catalyst is actually slightly more complicated than pure iron. It has potassium hydroxide added to it as a promoter - a substance that increases its efficiency.

The pressure The pressure varies from one manufacturing plant to another, but is always high. You can't go far wrong in an exam quoting 200 atmospheres. Recycling At each pass of the gases through the reactor, only about 15% of the nitrogen and hydrogen converts to ammonia. (This figure also varies from plant to plant.) By continual recycling of the unreacted nitrogen and hydrogen, the overall conversion is about 98%. Explaining the conditions The proportions of nitrogen and hydrogen The mixture of nitrogen and hydrogen going into the reactor is in the ratio of 1 volume of nitrogen to 3 volumes of hydrogen. Avogadro's Law says that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. That means that the gases are going into the reactor in the ratio of 1 molecule of nitrogen to 3 of hydrogen. That is the proportion demanded by the equation. In some reactions you might choose to use an excess of one of the reactants. You would do this if it is particularly important to use up as much as possible of the other reactant - if, for example, it was much more expensive. That doesn't apply in this case.

There is always a down-side to using anything other than the equation proportions. If you have an excess of one reactant there will be molecules passing through the reactor which can't possibly react because there isn't anything for them to react with. This wastes reactor space - particularly space on the surface of the catalyst. The temperature Equilibrium considerations You need to shift the position of the equilibrium as far as possible to the right in order to produce the maximum possible amount of ammonia in the equilibrium mixture. The forward reaction (the production of ammonia) is exothermic. According to Le Chatelier's Principle, this will be favoured if you lower the temperature. The system will respond by moving the position of equilibrium to counteract this - in other words by producing more heat. In order to get as much ammonia as possible in the equilibrium mixture, you need as low a temperature as possible. However, 400 - 450°C isn't a low temperature! Rate considerations The lower the temperature you use, the slower the reaction becomes. A manufacturer is trying to produce as much ammonia as possible per day. It makes no sense to try to achieve an equilibrium mixture which contains a very high proportion of ammonia if it takes several years for the reaction to reach that equilibrium.

You need the gases to reach equilibrium within the very short time that they will be in contact with the catalyst in the reactor. The compromise 400 - 450°C is a compromise temperature producing a reasonably high proportion of ammonia in the equilibrium mixture (even if it is only 15%), but in a very short time. The pressure Equilibrium considerations Notice that there are 4 molecules on the left-hand side of the equation, but only 2 on the right. According to Le Chatelier's Principle, if you increase the pressure the system will respond by favouring the reaction which produces fewer molecules. That will cause the pressure to fall again. In order to get as much ammonia as possible in the equilibrium mixture, you need as high a pressure as possible. 200 atmospheres is a high pressure, but not amazingly high. Rate considerations Increasing the pressure brings the molecules closer together. In this particular instance, it will increase their chances of hitting and sticking to the surface of the catalyst where they can react. The higher the pressure the better in terms of the rate of a gas reaction. Economic considerations Very high pressures are very expensive to produce on two counts. You have to build extremely strong pipes and containment vessels to withstand the very high pressure.

That increases your capital costs when the plant is built. High pressures cost a lot to produce and maintain. That means that the running costs of your plant are very high. The compromise 200 atmospheres is a compromise pressure chosen on economic grounds. If the pressure used is too high, the cost of generating it exceeds the price you can get for the extra ammonia produced. The catalyst Equilibrium considerations The catalyst has no effect whatsoever on the position of the equilibrium. Adding a catalyst doesn't produce any greater percentage of ammonia in the equilibrium mixture. Its only function is to speed up the reaction. Rate considerations In the absence of a catalyst the reaction is so slow that virtually no reaction happens in any sensible time. The catalyst ensures that the reaction is fast enough for a dynamic equilibrium to be set up within the very short time that the gases are actually in the reactor. Separating the ammonia When the gases leave the reactor they are hot and at a very high pressure. Ammonia is easily liquefied under pressure as long as it isn't too hot, and so the temperature of the mixture is lowered enough for the ammonia to turn to a liquid. The nitrogen and hydrogen remain as gases even under these high pressures, and can be recycled.

Questions to test your understanding If this is the first set of questions you have done, please read the introductory page before you start. You will need to use the BACK BUTTON on your browser to come back here afterwards. Case Study: The Haber Process - Chemistry LibreTexts Case Study: The Haber Process Last updated Contributors This page describes the Haber Process for the manufacture of ammonia from nitrogen and hydrogen, and then goes on to explain the reasons for the conditions used in the process. It looks at the effect of temperature, pressure and catalyst on the composition of the equilibrium mixture, the rate of the reaction and the economics of the process. A brief summary of the Haber Process The Haber Process combines nitrogen from the air with hydrogen derived mainly from natural gas (methane) into ammonia. The reaction is reversible and the production of ammonia is exothermic. \[ N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3 (g) \tag{ΔH=-92.4 kJ/mol}\] A flow scheme for the Haber Process looks like this: Some notes on the conditions The catalyst: The catalyst is actually slightly more complicated than pure iron.

It has potassium hydroxide added to it as a promoter - a substance that increases its efficiency. The pressure: The pressure varies from one manufacturing plant to another, but is always high. You can't go far wrong in an exam quoting 200 atmospheres. Recycling: At each pass of the gases through the reactor, only about 15% of the nitrogen and hydrogen converts to ammonia. (This figure also varies from plant to plant.) By continual recycling of the unreacted nitrogen and hydrogen, the overall conversion is about 98%. Explaining the conditions The proportions of nitrogen and hydrogen: The mixture of nitrogen and hydrogen going into the reactor is in the ratio of 1 volume of nitrogen to 3 volumes of hydrogen. Avogadro's Law says that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. That means that the gases are going into the reactor in the ratio of 1 molecule of nitrogen to 3 of hydrogen. That is the proportion demanded by the equation. In some reactions you might choose to use an excess of one of the reactants.

You would do this if it is particularly important to use up as much as possible of the other reactant - if, for example, it was much more expensive. That does not apply in this case. There is always a down-side to using anything other than the equation proportions. If you have an excess of one reactant there will be molecules passing through the reactor which can't possibly react because there isn't anything for them to react with. This wastes reactor space - particularly space on the surface of the catalyst. The temperature Equilibrium considerations: You need to shift the position of the equilibrium as far as possible to the right in order to produce the maximum possible amount of ammonia in the equilibrium mixture. The forward reaction (the production of ammonia) is exothermic. \[ N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3 (g) \tag{ΔH=-92.4 kJ/mol}\] According to Le Chatelier's Principle , this will be favored if you lower the temperature. The system will respond by moving the position of equilibrium to counteract this - in other words by producing more heat.

To get as much ammonia as possible in the equilibrium mixture, you need as low a temperature as possible. However, 400 - 450°C isn't a low temperature! Rate considerations: The lower the temperature you use, the slower the reaction becomes. A manufacturer is trying to produce as much ammonia as possible per day. It makes no sense to try to achieve an equilibrium mixture which contains a very high proportion of ammonia if it takes several years for the reaction to reach that equilibrium. You need the gases to reach equilibrium within the very short time that they will be in contact with the catalyst in the reactor. The compromise: 400 - 450°C is a compromise temperature producing a reasonably high proportion of ammonia in the equilibrium mixture (even if it is only 15%), but in a very short time. The pressure Equilibrium considerations \[ N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3 (g) \;\;\; \Delta H=-92 \, \text{kJ mol}^{-1}\] Notice that there are 4 molecules on the left-hand side of the equation, but only 2 on the right.

According to Le Chatelier's Principle, if you increase the pressure the system will respond by favoring the reaction which produces fewer molecules. That will cause the pressure to fall again. In order to get as much ammonia as possible in the equilibrium mixture, you need as high a pressure as possible. 200 atmospheres is a high pressure, but not amazingly high. Rate considerations: Increasing the pressure brings the molecules closer together. In this particular instance, it will increase their chances of hitting and sticking to the surface of the catalyst where they can react. The higher the pressure the better in terms of the rate of a gas reaction. Economic considerations: Very high pressures are very expensive to produce on two counts. You have to build extremely strong pipes and containment vessels to withstand the very high pressure. That increases your capital costs when the plant is built. High pressures cost a lot to produce and maintain. That means that the running costs of your plant are very high. The compromise: 200 atmospheres is a compromise pressure chosen on economic grounds. If the pressure used is too high, the cost of generating it exceeds the price you can get for the extra ammonia produced. The catalyst Equilibrium considerations: The catalyst has no effect whatsoever on the position of the equilibrium.

Adding a catalyst doesn't produce any greater percentage of ammonia in the equilibrium mixture. Its only function is to speed up the reaction. Rate considerations: In the absence of a catalyst the reaction is so slow that virtually no reaction happens in any sensible time. The catalyst ensures that the reaction is fast enough for a dynamic equilibrium to be set up within the very short time that the gases are actually in the reactor. Separating the ammonia: When the gases leave the reactor they are hot and at a very high pressure. Ammonia is easily liquefied under pressure as long as it isn't too hot, and so the temperature of the mixture is lowered enough for the ammonia to turn to a liquid. The nitrogen and hydrogen remain as gases even under these high pressures, and can be recycled. Nitric Acid By mixing one part ammonia to nine parts air with the use of a catalyst, the ammonia will get oxidized to nitric acid.

\[4 NH_3 + 5 O_2 \rightarrow 4 NO + 6 H_2O\] \[2 NO + O_2 \rightarrow 2 NO_2\] \[2 NO_2 + 2 H_2O \rightarrow 2 HNO_3 + H_2\] Contributors There are no recommended articles. Article type Haber Process for Ammonia Production Chemistry Tutorial (ii) very high pressure (&approx;250 atmospheres, &approx;25,500kPa) (iii) a catalyst (a porous iron catalyst prepared by reducing magnetite, Fe3O4). Osmium is a much better catalyst for the reaction but is very expensive. This process produces an ammonia, NH3(g), yield of approximately 10-20%. The Haber synthesis was developed into an industrial process by Carl Bosch. Please do not block ads on this website. No ads = no money for us = no free stuff for you!

Equilibrium Considerations The reaction between nitrogen gas and hydrogen gas to produce ammonia gas is an exothermic equilibrium reaction , releasing 92.4 kJ mol-1 of energy at 298 K (25oC). N2(g) ammonia + 92.4 kJ mol-1 By Le Chetalier's Principle , increasing the pressure on the reaction mixture favours the formation of ammonia gas: Increasing the pressure causes the equilibrium position to move to the right resulting in a higher yeild of ammonia since there are more gas molecules on the left hand side of the equation (4 in total) than there are on the right hand side of the equation (2). Increasing the pressure means the system adjusts to reduce the effect of the change, that is, to reduce the pressure by having fewer gas molecules. By Le Chetalier's Principle , decreasing the temperature of the reaction mixture favours the formation of ammonia gas: Decreasing the temperature causes the equilibrium position to move to the right resulting in a higher yield of ammonia since the reaction is exothermic (releases heat).

Reducing the temperature means the system will adjust to minimise the effect of the change, that is, it will produce more heat since energy is a product of the reaction, and will therefore produce more ammonia gas as well. However, the rate of the reaction at lower temperatures is extremely slow, so a higher temperature must be used to speed up the reaction which results in a lower yield of ammonia. The equilibrium expression for this reaction is: Kc = [NH3(g)]2 [N2(g)][H2(g)]3 As the temperature increases, the value of the equilibrium constant decreases as the yield of ammonia decreases. Temperature 1.5 x 10-5 lower Kc In summary, higher pressure favours formation of ammonia gas, lower temperature favours formation of ammonia gas BUT the reaction rate is slow at lower temperatures. Do you know this? Play the game now! Reaction Rate (Kinetic) Considerations A catalyst such as an iron catalyst is used to speed up the reaction by lowering the activation energy so that the N2 bonds and H2 bonds can be more readily broken.

Increased temperature means more reactant molecules have sufficient energy to overcome the energy barrier to reacting (activation energy) so the reaction is faster at higher temperatures (but the yield of ammonia is lower as discussed above). A temperature range of 400-500oC is a compromise designed to achieve an acceptable yield of ammonia (10-20%) within an acceptable time period. Refer to the graph on the right which records the yield of ammonia as a percentage at different pressures: At 200oC and pressures above 750 atm there is an almost 100% conversion of reactants to the ammonia product. Since there are difficulties associated with containing larger amounts of materials at this high pressure, lower pressures of around 200 atm are used industrially. By using a pressure of around 200 atm and a temperature of about 500oC, the yield of ammonia is 10-20%, while costs and safety concerns in the building and during operation of the plant are minimised During industrial production of ammonia, the reaction never reaches equilibrium as the gas mixture leaving the reactor is cooled to liquefy and remove the ammonia. The remaining mixture of reactant gases are recycled through the reactor.

The heat released by the reaction is removed and used to heat the incoming gas mixture. Do you understand this? urea Chemicals synthesis of: nitric acid, HNO3, which is used in making explosives such as TNT (2,4,6-trinitrotoluene), nitroglycerine which is also used as a vasodilator (a substance that dilates blood vessels) and PETN (pentaerythritol nitrate).

sodium hydrogen carbonate (sodium bicarbonate), NaHCO3 sodium carbonate, Na2CO3 hydrogen cyanide (hydrocyanic acid), HCN hydrazine, N2H4 (used in rocket propulsion systems) Explosives ammonium nitrate, NH4NO3 Fibres and Plastics nylon , -[(CH2)4-CO-NH-(CH2)6-NH-CO]-,and other polyamides Refrigeration used for making ice, large scale refrigeration plants, air-conditioning units in buildings and plants Pharmaceuticals used in the manufacture of drugs such as sulfonamide which inhibit the growth and multiplication of bacteria that require p-aminobenzoic acid (PABA) for the biosynthesis of folic acids, anti-malarials and vitamins such as the B vitamins nicotinamide (niacinamide) and thiamine. Pulp and Paper ammonium hydrogen sulfite, NH4HSO3, enables some hardwoods to be used Mining and Metallurgy used in nitriding (bright annealing) steel, used in zinc and nickel extraction Cleaning ammonia in solution is used as a cleaning agent such as in 'cloudy ammonia' Can you apply this? Take the exam now!

A Brief History of Ammonia Production At the beginning of the 20th century there was a shortage of naturally occurring, nitrogen-rich fertilisers, such as Chile saltpetre, which prompted the German Chemist Fritz Haber, and others, to look for ways of combining the nitrogen in the air with hydrogen to form ammonia, which is a convenient starting point in the manufacture of fertilisers. This process was also of interest to the German chemical industry as Germany was preparing for World War I and nitrogen compounds were needed for explosives. The hydrogen for the ammonia synthesis was made by the water-gas process (a Carl Bosch invention) which involves blowing steam through a bed of red hot coke resulting in the separation of hydrogen from oxygen. The nitrogen was obtained by distillation of liquid air, then by cooling and compressing air.

These days, the hydrogen is produced by reforming light petroleum fractions or natural gas (methane, CH4) by adding steam: CH4(g) + H2O(g) The Haber Process | Grade 12U Chemistry-Systems and Equilibrium Grade 12U Chemistry-Systems and Equilibrium The Haber process, this process is also referred to as Haber – Bosch process This Process is basically the reaction of nitrogen and hydrogen to produce ammonia So how does this process happen, the steps of this process? The nitrogen gas (N2) is reacted with the hydrogen gas (H2) over an iron catalyst (fe3+) and Aluminum oxide (Al2O3) and Potassium oxide (k2O) are used as products. Then the unreacted gases are sent back in to the reactor and create ammonia. And the gases that have reacted are cooled and ammonia turns to liquid. The Haber process Explaing the pic above: Why is the gas recycled?

The unreacted gas is recycled because each time the gas passes through the reactor only about 15% of the hydrogen and nitrogen convert to ammonia so the other 85% of the gas is recycled and reused and are reacted over the iron catalyst (fe3+) once again. The reason potassium oxide and aluminum oxide are used as a product alongside with the iron catalyst is to increases the efficiency of the process. Avigardro’s law: This law states that “equal volumes of gases at the same temp and pressure contain equal numbers of molecules” because of this law we know that the ratio for nitrogen and hydrogen is 1:3. Equilibruim: Equation: This equation must be shifted to as far as possible to the right to produce the most ammonia. Since the reaction is shifted to the right is a forward reaction and it is exothermic According to Le Chatelier’s principle: in order to get the equation to shift as far as possible to the right a low temperature is needed. Since this reaction is exothermic and the temperature is decreased the yield of the products is then increased so this will make more ammonia.

And since the lowest temperature possible is 400 to 450°C, this is one of the reasons why only about 15% of the hydrogen and nitrogen react and produce ammonia. The reason why 400 to 450°C is because if the temperature of the reaction is low the reaction will slow down and it would take longer to produce ammonia which in the long run is uneconomical. So the temperature 400 to 450°C is a compromise temperature where the yield is high enough to get a reasonable rate of reaction and a good yield of ammonia. Shown in the equation above there are 4 molecules in total on the left side and only 2 molecules on the right side. And according to Le Chateliers principle it states that if the pressure of the system is increased the system will respond by favoring the side of the reaction that produces or has less molecules so in this cases since there are only 2 molecules on the right side and 4 on the left side, the forward reaction is favored and ammonia will be produced. The reason why a high pressure is needed is because it brings the molecules together and this increases their chance of hitting each other and sticking to the catalyst so they can react.

The catalyst itself has no effect on the equilibrium; the catalyst is only there to speed up the time needed for this reaction. But the catalyst does ensure that the reaction is fast enough for a dynamic equilibrium can be set up while both the gases are in the reactor. Effect On us: The Haber process has good effects such as it creates fertilizer that feeds more than one half of the world. On the bad side the Haber process has bad effects on the soil, when the fertilizer is added to the soil, the soil has gained a lot more of nitrogen due to this the organic matter in the soil is lost. Plants only absorb about 10% nutrients and about 50% of the fertilizer applied is actually in the crop that it was used on. So there is a lot of wasted fertilizer and lastly a lot of the Haber process nitrogen is used in making bombs and bullets for guns. Many of the road side bombs in many parts of the world are made by using the Haber process. Uses of Haber process: During World War 1 the Haber process was used for the production of nitric acid. The Haber process is still very useful today since ammonia is needed for fertilizer.

And ammonia also makes cleaning supplies that are used to clean glass objects. Which the help of The Haber process roughly about 500 million tons of fertilizer is produced in a year, And this fertilizer feeds about 40% of the world’s population. Refrences: Ammonia Ammonia Ammonia The manufacture of ammonia is crucial for the world's agricultural industry for from it all fertilizers that contain nitrogen are produced. Uses of ammonia The manufacture of fertilizers is by far the most important use of ammonia. These include urea , ammonium salts (ammonium phosphates, ammonium nitrate, calcium ammonium nitrate) and solutions of ammonia. Figure 1 The uses of ammonia. An increasing amount of ammonia, although still small compared with other uses, is used as a concentrated solution in combating the discharge of nitrogen oxides from power stations. Annual production of ammonia Ammonia ranks second, to sulfuric acid, as the chemical with the largest tonnage. It is being increasingly made in countries which have low cost sources of natural gas and coal (China and Russia account for ca 40%).

The largest plants produce about 3000 tonnes a day and there are plans to build plants that produce 4000-5000 tonnes a day, which would mean that the total world output could be managed with 100 such units. Current production is: World Data from: 1. Federal State Statistics Service: Russian Federation 2011 The world population is increasing by about 1.4% a year and the increase in tonnage of ammonia made just about keeps pace. Manufacture of ammonia The manufacture of ammonia from nitrogen and hydrogen takes place in two main stages: a) the manufacture of hydrogen b) the synthesis of ammonia (the Haber Process) The manufacture of hydrogen involves several distinct processes. Figure 2 shows their sequence and the location within an ammonia plant (steps1-5). The converter used to make ammonia from the hydrogen is also shown (step 6). What occurs in each of these steps is described below the figure. Figure 2 An ammonia plant in Western Australia: 1 Desulfurisation units (a) The manufacture of hydrogen Hydrogen is produced from a variety of feedstocks, mostly from natural gas, coal or naphtha.

The ways in which hydrogen is obtained from these feedstocks are dealt with separately. Hydrogen from natural gas (methane) This involves two stages: i) the manufacture of synthesis gas (a mixture of carbon monoxide and hydrogen (steam reforming)) >ii) the removal of the carbon monoxide and production of a mixture of hydrogen and nitrogen (the shift reaction) (i) The manufacture of synthesis gas Whichever way the methane is obtained, it will contain some organic sulfur compounds and hydrogen sulfide, both of which must be removed. Otherwise, they will poison the catalyst needed in the manufacture of synthesis gas. In the desulfurisation unit, the organic sulfur compounds are often first converted into hydrogen sulfide, prior to reaction with zinc oxide. The feedstock is mixed with hydrogen and passed over a catalyst of mixed oxides of cobalt and molybdenum on an inert support (a specially treated alumina) at ca 700 K.

Then the gases are passed over zinc oxide at ca 700 K and hydrogen sulfide is removed: Primary steam reforming converts methane and steam to synthesis gas, a mixture of carbon monoxide and hydrogen: High temperatures and low pressures favour the formation of the products (Le Chatelier's Principle). In practice, the reactants are passed over a catalyst of nickel, finely divided on the surface of a calcium oxide/aluminium oxide support contained in vertical nickel alloy tubes. The tubes, up to 350 in parallel, are heated in a furnace above 1000 K and under a pressure of ca 30 atm. This is an example of a tubular reactor . Secondary steam reforming reacts oxygen from the air with some of the hydrogen present and the resulting mixture is passed over a nickel catalyst. The steam and heat produced from the combustion reforms most of the residual methane. Among the key reactions are: The emerging gas from this net exothermic stage is at ca 1200 K and is cooled in heat exchangers . The steam formed from the water used in cooling the gases is used to operate turbines and thus compressors and to preheat reactants.

Some recent designs use waste heat from the secondary reformer directly to provide heat for the primary reformer. At this stage the gas contains hydrogen, nitrogen, carbon monoxide and carbon dioxide and about 0.25% methane. As air contains 1% argon, this also accumulates in the synthesis gas. (ii) The shift reaction This process converts carbon monoxide to carbon dioxide, while generating more hydrogen. It takes place in two stages. In the first, the high temperature shift reaction, the gas is mixed with steam and passed over an iron/chromium(III) oxide catalyst at ca 700 K in a fixed bed reactor . This decreases the carbon monoxide concentration from 11%: In the second stage, the low temperature shift reaction, the mixture of gases is passed over a copper-zinc catalyst at ca 500 K. The carbon monoxide concentration is further reduced to 0.2%. The reaction is done in two stages for several reasons. The reaction is exothermic. However, at high temperature, the exit concentration of carbon monoxide is still quite high, due to equilibrium control.

The copper catalyst used in the low temperature stage is very sensitive to high temperatures, and could not operate effectively in the high temperature stage. Thus, the bulk of the reaction is carried out at high temperature to recover most of the heat. The gas is then removed at low temperature, where the equilibrium is much more favourable, on the very active but unstable copper catalyst. The gas mixture now contains about 18% carbon dioxide which is removed by scrubbing the gas with a solution of a base, using one of several available methods. One that is favoured is an organic base (in the carbon dioxide absorber), a solution of an ethanolamine, often 2,2'-(methylimino)bis-ethanol (N-methyl diethanolamine). The carbon dioxide is released on heating the solution in the carbon dioxide stripper). Much of it is liquefied and sold, for example, for carbonated drinks, as a coolant for nuclear power stations and for promoting the growth of plants in greenhouses . The last traces of oxides of carbon are removed by passing the gases over a nickel catalyst at 600 K: This process is known as methanation.

A gas is obtained of typical composition: 74% hydrogen, 25% nitrogen, 1% methane, together with some argon. Hydrogen from naphtha If naphtha is used as the feedstock, an extra reforming stage is needed. The naphtha is heated to form a vapour, mixed with steam and passed through tubes, heated at 750 K and packed with a catalyst, nickel supported on a mixture of aluminium and magnesium oxides. The main product is methane together with oxides of carbon, and is then processed by steam reforming, as if it was natural gas, followed by the shift reaction. Hydrogen from coal If coal is used, it is first finely ground and heated in an atmosphere of oxygen and steam. Some of the coal burns very rapidly in oxygen (in less than 0.1 s) causing the temperature in the furnace to rise and the rest of the coal reacts with the steam: The gas emitted contains ca 55% carbon monoxide, 30% hydrogen, 10% carbon dioxide and small amounts of methane and other hydrocarbons. This mixture is treated by the shift reaction.

The main problems of using coal includes the large amounts of sulfur dioxide and trioxide generated in burning coal and the significant amounts of other impurities such as arsenic and bromine, all of which are very harmful to the atmosphere and all of which are severe poisons to the catalysts in the process. There is also a massive problem with disposal of the ash. (b) The manufacture of ammonia (The Haber Process) The heart of the process is the reaction between hydrogen and nitrogen in a fixed bed reactor . The gases, in stoichiometric proportions, are heated and passed under pressure over a catalyst (Figure 3). Figure 3 A diagram illustrating a conventional synthesis reactor (a converter). The proportion of ammonia in the equilibrium mixture increases with increasing pressure and with falling temperature (Le Chatelier's Principle). Quantitative data are given in Table 1. To obtain a reasonable yield and favourable rate, high pressures, moderate temperatures and a catalyst are used. Table 1 Percentage, by volume, of ammonia in the equilibrium mixture for the reaction between nitrogen and hydrogen at a range of temperatures and pressures. A wide range of conditions are used, depending on the construction of the reactor.

Temperatures used vary between 600 and 700 K, and pressures between 100 and 200 atmospheres. Much work is being done to improve the effectiveness of the catalyst so that pressures as low as 50 atmospheres can be used. As the reaction is exothermic, cool reactants (nitrogen and hydrogen) are added to reduce the temperature of the reactors (Figure 3). The ammonia is usually stored on site (step 7) and pumped to another part of the plant where it is converted into a fertilizer (urea or an ammonium salt). However it is sometimes transported by sea (Figure 4) or by road, to be used in another plant. Figure 4 In a plant in Western Australia, the ammonia is transferred by pipeline to a nearby harbour (Figure 2, step 8) and transported by ship. This one is carrying about 40 000 tonnes of liquefied ammonia. By kind permission of Yara International ASA. The original catalyst that Haber used was Fe3O4, which was reduced by the reactant, hydrogen, to iron.

Much work was done to improve the catalyst and it was found that a small amount of potassium hydroxide was effective as a promoter . Recently research has been focussed on finding even more effective catalysts to enable the process to take place at lower pressures and temperatures. Ruthenium on a graphite surface is a promising one. Figure 5 The production of ammonia. Postscript The Haber Process is of such importance to our lives that it has figured in three Nobel Prizes in chemistry, all to German scientists, over a period of nearly 90 years, a remarkable record. The first was given in 1918, to Fritz Haber, the chemist who developed the process in the laboratory. The second was to Carl Bosch, whose brilliant engineering skills made the process viable on a massive scale, but who waited until 1931 for his award. In 2007, Gerhard Ertl was awarded the Prize for his work on catalysis of gaseous reactions on solids . Among the wide range of reactions he studied, he gained evidence for the adsorption of nitrogen molecules and hydrogen molecules on the surface iron and that these adsorbed molecules dissociate into atoms.

These atoms then join up in stages to form the ammonia molecule. It must be remembered that the conditions used in these studies (at less than 10-10 atm) are very different from the conditions used in industry, ca 150 atm. Summary The whole process of producing ammonia from methane is summarized in Figure 5. If coal or naphtha is the feedstock, extra processes are needed. Naphtha is converted into methane and oxides of carbon before going into the primary reformer and thence to the shift reaction. Coal is also converted into hydrogen and carbon oxides and this mixture then undergoes the shift reaction. BBC - GCSE Bitesize: Ammonia and the Haber process Ammonia and the Haber process Ammonia Ammonia (NH3) is a compound [compound: A compound is a substance formed by the chemical union (involving bond formation) of two or more elements. ] of nitrogen and hydrogen. It is a colourless gas with a choking smell, and a weak alkali which is very soluble in water. Ammonia is used to make fertilisers, explosives, dyes, household cleaners and nylon.

It is also the most important raw material in the manufacture of nitric acid. Ammonia is manufactured by combining nitrogen and hydrogen in an important industrial process called the Haber process. An ammonia production plant. Photo courtesy of WMC Resources Ltd The Haber process The raw materials for this process are hydrogen and nitrogen. Hydrogen is obtained by reacting natural gas - methane - with steam, or through the cracking of oil. Nitrogen is obtained by burning hydrogen in air. Air is 80 per cent nitrogen; nearly all the rest is oxygen. When hydrogen is burned in air, the oxygen combines with the hydrogen, leaving nitrogen behind. Nitrogen and hydrogen will react together under these conditions: a high temperature - about 450ºC a high pressure - about 200 atmospheres (200 times normal pressure) an iron catalyst The Haber Process - Nitrogen Fertilizer from the Air Putting the Carbon Back in your Soil The Haber Process Nitrogen Fertilizer from the Air The Haber Process, also called the Haber-Bosch Process, is a complex chemical procedure that takes nitrogen from the air and under high pressures and temperatures combines it with hydrogen to produce ammonia.

This ammonia is the base of the synthetic nitrogen fertilizers increasingly used around the world today. Nitrogen, a key component of all proteins, DNA, and RNA, is vital to life here. Plants can only use fixed nitrogen and the lack of fixed nitrogen is often the limiting factor in an ecosystem or for crops. Our ability to fix nitrogen ourselves seemed to be a perfect solution. However, the enormous blessings of the Haber Process are balanced by some pretty serious curses. On the blessing side synthetic nitrogen fertilizer produced by the Haber Process is credited with feeding a third to half the present world population. In fact about half the nitrogen in each of our bodies is there thanks to the Haber Process. On the curse side we have several issues including: Serious imbalances to the nitrogen cycle. High fossil fuel energy inputs. Negative effects on soil organisms and soil organic matter. Excess runoff cause ocean dead zones. Major component of weapons including all those roadside bombs. A Very Short History Through the 1800's, heavily populated countries in Europe such as Germany and Britain used guano from islands off the coast of Peru and salt petre from Chile as a source of natural nitrogen fertilizer.

This trade was much like our current oil trade in importance both for food security and politically. By the end of the 19th century these sources of fixed nitrogen were running out. Scientists were challenged to come up with a process of making fixed nitrogen from the nitrogen in the air. The winner of the race was Fritz Haber. His tabletop contraption used high pressure, heat and a catalyst to slowly, drop by drop, synthesize a cup of ammonia over the course of two hours. The German chemical company BASF bought the process and Carl Bosch was charged with upscaling it to factory size, a feat he managed by 1913. The process, often called the Haber Bosch Process requires high pressures, 200 - 300 atmospheres, and high temperatures, 400 to 550 degrees C. For a brief time the first factory turned out fertilizer. With onset of WW1 Germany needed munitions for its war effort. The factory was seconded to munitions and is credited or blamed with greatly prolonging WW1 and making WW2 possible. These synthetic nitrogen fertilizers are explosive.

This photo shows the aftermath of the 1921 Oppau explosion that took place at the BASF ammonium nitrate factory there.Between 500 and 600 were killed and 2000 injured. Impacts of the Haber Process on Farming, Gardening and the World More Food - More People It wasn't until the 1950's, after the first and second world wars, that the Haber process really started to affect farming. Ammonia stocks, diverted in wartime to make bombs and bullets, started being used to produce the synthetic nitrogen fertilizers used everywhere today. Increased food production has allowed the human population to swell to its current almost 7 billion size. Estimates are that a third of the earth's human population is fed thanks to the Haber process. Many people believe that to stop or limit the use of synthetic fertilizers would lead to mass starvation. Balance in the Nitrogen Cycle Our impact on the nitrogen cycle is actually bigger than our effect on the carbon cycle. We convert more N2 gas into fixed reactive forms than all the Earth's processes combined. Synthetic nitrogen fertilizer production, vehicle exhaust emissions and even growing nitrogen fixing crops all fix nitrogen.

We add an estimated 140 million tons of fixed nitrogen of which roughly 80-100 million tons is from the Haber process chemical plants. If you study scientific material on the nitrogen cycle you can't help but notice that molecules of nitrogen that get fixed can also get unfixed. Nitrogen fixing bacteria are balanced with denitrifying counterparts in roughly equal proportions. Over time there might be a slow increase in fixed nitrogen but in the natural world the shift is very slow. The Haber process contributions of reactive nitrogen are not offset by denitrification. A group of scientists led by Johan Rockstrom explored the questions of planetary boundaries in an article recently published by Nature. The usual climate change issue was of course reported. What might surprise you though is that while we have only just slipped past the edge of the CO2 boundary, the nitrogen cycle disruption was deemed to be about three times over the safe limit. This graphic illustrates the extent to which we may be overstepping the nitrogen boundary. These scientists suggest we reduce our fixed nitrogen additions to about 35 million tons or roughly a third of what we are currently slinging onto fields and belching out of cars and factories.

Haber Process is an Energy Glutton Given that the Haber process requires temperatures of 400 - 550C and pressures of 200 - 300 atmospheres it's not surprising that it uses a lot of energy. Manufacture of nitrogen fertilizers uses about 5% of the world's natural gas production, equivalent to 1-2% of the world's annual energy consumption. While natural gas is among the more plentiful fossil fuels, this level of use is not sustainable in the long term. As well, as oil reserves worldwide dwindle, prices will rise putting fertilizers some farmers have come to depend on out of reach for the poor. Negative Effects on Soil and Soil Community A balanced diet for soil organisms in simple terms is about 20 parts carbon for 1 part nitrogen. Adding nitrogen fertilizer to soil usually creates a surplus of nitrogen. To balance their diet the soil microbes go after carbon reserves in the soil - first the easily digested stuff and then the humus. The organic matter in the soil is lost. The soil food web, that intricate web of civilization living in the soil, is starved and begins to die.

The functions they performed, whether they be fixing nitrogen, making phosphorus stores available or defending crops against various pests and diseases, are also lost. As the natural system breaks down the soil organisms who survive tend to be the bad guys. Their usual enemies are gone. Erosion of topsoil increases; pests and diseases increase. Not a good scenario. Farmers, to maintain yields, turn to more chemical solutions. More fertilizers, more herbicides and pesticides... and more trouble for the community making it's home in the dirt. As well, recent research at the Center for Ecology and Evolutionary Biology at the University of Oregon by Dr. Jennifer Fox, found that communication links between plants and nitrogen fixing bacteria are damaged by use of excessive synthetic nitrogen fertilizer. This leads to less nitrogen fixed naturally and increased reliance on synthetics. Fertilizer Runoff and Pollution Plants can absorb the nutrients available through chemical fertilizers. However, often the dose is too much too fast, and much of the fertilizer is wasted. As little as 10% may be actually absorbed by the crop and rarely more than 50% of the fertilizer applied makes it into the crop grown. The nitrogen fertilizer tends to be especially mobile in the soil.

It passes right by the plants and off into ground water polluting wells or is eroded into surface streams. The huge ocean dead zones are a result of nitrogen and other fertilizers and farm chemicals ending up in watersheds bound for the oceans. Bombs and Bullets There is still a good chunk of Haber process nitrogen that winds up in war zones. In fact many of those roadside bombs terrorizing people in various parts of the world are made in part from Nitrogen fertilizer synthesized with this process. Hope - Compost and Organic Matter Restores Soils The synthetic nitrogen fertilizers brought to us complements of the Haber process seemed at first a boon to humanity. Synthetic chemical fertilizers have only been commonly available in the last 50 years. Using synthetic fertilizers has become so routine farmers who use them are considered conventional while those using the techniques thousands of years old are called organic. Given that so little of the Haber process fertilizer actually makes it into crops (and so much into the ocean) it seems to me that we may be able to feed the world better with a lot less chemical additions if we pay a lot more attention to biology. On a personal scale making and using compost is a step that restores soil and their communities.

Can that be a bad thing? Want more information? These books will let you delve more deeply into this topic. The Alchemy of Air is Fritz Haber's biography. I'd balance it with one or more of the other books to explore the hope of soil biology. Canadians can access the Amazon.ca store from the second set of links. Check Out These Pages Haber-Bosch Process - Chemistry Information The Haber process or Haber-Bosch process is the primary industrial method used to make ammonia or fix nitrogen . The Haber process reacts nitrogen and hydrogen gas to form ammonia: N2 + 3 H2 → 2 NH3 (ΔH = −92.4 kJ·mol−1) History of the Haber Process Fritz Haber , a German chemist, and Robert Le Rossignol, a British chemist, demonstrated the first ammonia synthesis process in 1909. They formed ammonia drop by drop from pressurized air. However, the technology did not exist to extend the pressure required in this tabletop apparatus to commercial production. Carl Bosch , an engineer at BASF, resolved the engineering problems associated with industrial ammonia production. BASF's German Oppau plant commenced ammonia production in 1913.

How the Haber-Bosch Process Works Haber's original process made ammonia from air. The industrial Haber-Bosch process mixes nitrogen gas and hydrogen gas in a pressure vessel that contains a special catalyst to speed the reaction. continue reading below our video What are the Seven Wonders of the World From a thermodynamic standpoint, the reaction between nitrogen and hydrogen favors the product at room temperature and pressure, but the reaction does not generate much ammonia. The reaction is exothermic; at increased temperature and atmospheric pressure, the equilibrium quickly switches the other direction. So, the catalyst and increased pressure are the scientific magic behind the process. Bosch's original catalyst was osmium, but BASF quickly settled upon a less-expensive iron-based catalyst, which is still in use today. Some modern processes employ a ruthenium catalyst, which is more active than the iron catalyst. Although Bosch originally electrolyzed water to obtain hydrogen, the modern version of the process uses natural gas to obtain methane, which is processed to get hydrogen gas. It is estimated that 3-5% of the world's natural gas production goes toward the Haber process. The gases pass over the catalyst bed multiple times, since conversion to ammonia is only around 15% each time.

By the end of the process, about 97% conversion of nitrogen and hydrogen to ammonia is achieved. Importance of the Haber Process Some people consider the Haber process to be the most important invention of the past 200 years! The primary reason the Haber process is important is because ammonia is used as a plant fertilizer, enabling farmers to grow enough crops to support an ever-increasing world population. The Haber process supplies 500 million tons (453 billion kilograms) of nitrogen-based fertilizer annually, which is estimated to support food for a third of the people on Earth. There are negative associations with the Haber process, too. In World War I, the ammonia was used to produce nitric acid to manufacture munitions. Some argue the population explosion, for better or worse, would not have happened without the increased food available because of the fertilizer. Also, the release of nitrogen compounds has had a negative environmental impact. References Enriching the Earth: Fritz Haber, Carl Bosch, and the Transformation of World Food Production, Vaclav Smil (2001) ISBN 0-262-19449-X.

US Enivronmental Protection Agency: Human Alteration of the Global Nitrogen Cycle: Causes and Consequences by Peter M. Vitousek, Chair, John Aber, Robert W. Howarth, Gene E. Likens, Pamela A. Matson, David W. Schindler, William H. Schlesinger, and G. David Tilman BBC - GCSE Bitesize Science - Ammonia : Revision Next The Haber process uses a reversible reaction between nitrogen and hydrogen to form ammonia. This is used to make nitrogenous fertilisers which promote plant growth. If overused, fertilisers can pollute rivers and lakes and cause ‘eutrophication’ - algae overgrowth which prevents sunlight reaching other water plants and causes the river or lake to eventually become lifeless. Ammonia and the Haber process Ammonia (NH3) is a compound [compound: A substance formed by the chemical union (involving bond formation) of two or more elements.] of nitrogen and hydrogen.

It is a colourless gas with a choking smell, and a weak alkali which is very soluble [soluble: Able to dissolve (usually in water).] in water It is used to make fertilisers, explosives, dyes, household cleaners and nylon. And it is also the most important raw material in the manufacture of nitric acid. Ammonia is manufactured by combining nitrogen and hydrogen in an important industrial process called the Haber process. An industrial ammonia plant The Haber process The raw materials for this process are hydrogen and nitrogen. Hydrogen is obtained by reacting natural gas - methane - with steam, or through the cracking [cracked: The breaking down of large hydrocarbon molecules into smaller, more useful hydrocarbon molecules by vapourizing them and passing them over a hot catalyst.] of oil. Nitrogen is obtained from air. Nitrogen and hydrogen will react together under these conditions: a high temperature - about 450ºC a high pressure - about 200 atmospheres (200 times normal pressure) an iron catalyst [catalyst: A catalyst changes the rate of a chemical reaction without being changed by the reaction itself.] The reaction is reversible.

N2(g) + 3H2(g) ⇌ 2NH3(g) The (g) indicates that the substance is a gas. The flow chart below shows the main stages in the Haber process. The reaction is reversible, and some nitrogen and hydrogen remain mixed with the ammonia. The reaction mixture is cooled so that the ammonia liquefies and can be removed. The remaining nitrogen and hydrogen are recycled. The Haber process for making ammonia The formation of ammonia in the Haber process In order to see this content you need to have both Javascript enabled and Flash installed. Page

Question: The 'Haber Process' is a reaction of nitrogen and hydrogen in order to produce which gas?
Answer: AMMONIA